Bonding (hl)
The actual arrangement of the atoms around the central atom in a molecule.
In order to find this we need to consider the number of regions of electron
density (electron pairs or unpaired single electrons in some cases) around
the central atom. The electrons will repel as far apart as possible as
they have the same charge.
The valence shell electron pair repulsion theory (VSEPRT)
gives us a means of working out the shapes of molecules and ions.
-
Draw the Lewis structure to show all of the valence shell electrons.
-
Then count the number of regions of electron density - this can
then be translated into the electronic shape.
-
Now consider the number of attached atoms and their orientation
keeping the lone (non-bonding) pairs as far apart as possible without
changing the electronic shape.
-
The resultant molecular arrangement gives the molecular shape. Note
that it may be slightly distorted by repulsions between regions of
electron density.
For a fuller account see VSEPRT
Molecules in three dimensions
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and when the viewing screen opens right click on the image to see the
options.
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here)
Predicting
molecular shapes 
Summary of the possible geometries
|
No. of electron pairs around the central atom
|
electronic shape
|
no of attached atoms
|
molecular shape
|
|
2
|
linear |
2
|
linear |
|
3
|
trigonal planar |
3
|
trigonal planar |
|
4
|
tetrahedral |
4
|
tetrahedral |
|
4
|
tetrahedral |
3
|
pyramidal |
|
4
|
tetrahedral |
2
|
angular or bent |
|
5
|
trigonal bipyramidal |
5
|
trigonal bipyramidal |
|
5
|
trigonal bipyramidal |
4
|
saw-horse (seesaw) |
|
5
|
trigonal bipyramidal |
3
|
T-shaped |
|
5
|
trigonal bipyramidal |
2
|
linear |
|
6
|
octahedral |
6
|
octahedral |
|
6
|
octahedral |
5
|
umbrella |
|
6
|
octahedral |
4
|
square planar |
This model explains the tetrahedral geometry of carbon and other atoms.
The electron structure of carbon is 1s2 2s2 2p2
suggesting that it should only be able to form two bonds (using the two
singly occupied orbitals). However it is known to make four single bonds
in many compounds and indeed never forms just two bonds. This can be explained
by hybridisation - the mixing of atomic orbitals producing degenerate
orbitals used for bonding.
-
sp3 hybridisation occurs when the 2s and 2p orbitals
merge to become sp3 orbitals (all of equal energy, length
etc.).
-
sp2 is the same except only two of the p orbitals are
hybridised, leaving one p orbital unchanged
-
sp is the same except only one of the p orbitals is hybridised and
two p orbitals are left unchanged
| hybridisation |
geometry |
example |
carbon in |
| sp3 |
tetrahedral |
methane |
CH4 |
| sp2 |
trigonal planar |
ethene |
CH2=CH2 |
| sp |
linear |
ethyne |
CH=CH |
When a particular molecule can be represented as several different Lewis
structures is is generally not actually any of these, but a hybrid (mixture)
of all of them. This can be represented either by using delocalised electrons,
or through resonance (where each possible structure is drawn and the actual
state 'resonates' between them. The delocalisation of these pi electrons
(which is effectively what happens) makes the molecule more stable (as
evidenced by lower energy) and gives the bonds a shorter length than would
be expected.
Examples:
|
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